Download HChem Unit 3 - Chpt11 Modern atom

Unit 3 - The Modern Atom
• What is our model of the Atom?
• What is wrong with it?
•
Homework: pg. 333-336 Q&P # 7, 8, 12-14,
20, 25, 31, 32, 36-39, 45, 50, 59, 70, 73, 75,
80, 81
Electromagnetic Spectrum
review
• What are some parts of the spectrum?
• It propagates how?
• What are the parts of a wave? how are
they related?
• frequency, nu ν
• wavelength, lambda
λ
• c = λ x ν c = speed of light
• photons - packets of light energy
Electromagnetic Radiation
 (lamba) = wavelength (m)
 (nu) = frequency (Hertz, Hz or s-1)
E = energy
c = speed of light, 2.9979 x 108 m/s
c =   they are inversely related
Know the relative order of radiation in E,   
EM Spectrum
EM spectrum visual
Emission from atoms
• Atoms absorb or emit energy.
• Becomes “excited” when absorbs
energy
• Can release energy by emitting a
photon.
Hydrogen
Helium
Carbon
Energy Levels of H atom
• The H atom has the experimental
spectrum from the previous slide
• Only certain types of photons were
produced!
• The H atom must have certain discrete
energy levels. Levels are quantized.
• Model constructed by Bohr - like
planets orbiting the sun.
• E=hxν
h = planck’s constant
Birth of Quantum Mechanics
• Black-body radiation - why does piece of iron
change colors when heated?
• Planck’s explanation
E=hxν
• Photoelectric effect - light shining on a metal
surface can emit electrons from surface.
• Einstein - light behaves like particle
• Bohr’s explanation of H-atom • de Broglie - wavelength of a particle!!!! little book pg 139
144
little book pg
Photoelectric Effect
Light with frequency lower than a
specific threshold have no electrons
emitted (no matter how intense it is)
Light with frequency greater than
threshold emits electrons and number
of electrons increases with intensity
Where is this going?
• Classical mechanics - visible objects at
ordinary velocities
• Quantum mechanics - describes behavior of
extremely small objects at velocities near the
speed of light - (Dr. Quantum video)
• Heisenberg - Uncertainty Principle impossible to know both position and
momentum of object at same time - little book pg 148
(slit video)
Wave Mechanical Model
• Schrodinger solved wave problem
mathematically ! - no physical meaning
• Orbits became orbitals - a region of
space with a probability of finding the
electron.
• Note:
calculate the max probability for
the H atom - you get 53pm, which is
the same as calculated from Bohr’s
model.
Where are the electrons?
• These energy level rules tell us where the electrons
are and how they are arranged in an atom.
• Principle Energy level n= 1,2,3,4,...
• Each Principle energy level has sublevels (orbitals)
• n=1 has 1 sublevel, n=2 has 2 sublevels, n=3 has 3
sublevels
• types of orbitals s, p, d, and f
• s has 1 orbital, p has 3 orbitals, d has 5 orbitals, f has 7
orbitals
• Now what?
12
Energy Levels of orbitals
As we keep adding energy levels,
we see as the principle quantum
number, n, increases the number
of sublevels (types of orbitals)
increases. In addition the energy
spacings get closer together 1s 2s - 3s - 4s - etc. So the energy of
the 4s orbital comes lower than the
3d. The order need not be
memorized because the elements
in Periodic Table shows it with its
s,p,d,f blocks.
Shapes of p and d orbitals
Electron Configuration rules
1. Electron’s occupy lowest energy level first aufbau principle
2. Maximum of 2 electrons in any orbital - Pauli
exclusion principle
– If 2 electrons occupy the same orbital they have
opposite spins. +1/2 or -1/2 also called spin up /
down or clockwise / counter-clockwise
3. For degenerate orbitals (the same energy like
the three p, five d, or seven f) use Hund’s rule,
also known as the bus rule - only pair up the
electrons if necessary.
Periodic Table
• Practice some electron configurations
–
–
–
–
–
–
H 1electron - 1s1
He 2 electrons - 1s2
2
1
Li 3 electrons - 1s 2s
Be - 1s22s2
B - 1s22s22p1 write e- configurations up to Ar
filling p orbitals - remember bus rule (unpaired e-)
also p has 3 orbitals so 6 total e- can fit in them
• Valence electrons (outermost e-) -largest n
– electron configuration of just valence e– electron dot notation
General s,p,d,f blocks
The periodic table
clearly shows that
after the 3p orbital,
the 4s fills before
the 3d. Likewise, 6s
4f 5d 6p is the order
when the
lanthanides start.
Periodic Table
7 Periods and 18 Groups
Group 1 - alkali metals - most reactive metals, Fr
most
Group 2 - alkali earth metals - next most reactive
Group 17 - halogens - most reactive nonmetals, F
most
Group 18 - noble gases - colorless, odorless,
nonreactive
Valence electrons, normal ionic charge (oxidation
number)
gain or lose electrons when make ionic compounds
electron configurations
Periodic Properties
Most properties of elements follow a trend in the
periodic table.
2 types of trends
Group trend - properties change some as the
elements go up or down a column.
Period trend - properties change some as the
elements go across a row.
Ionization Energy
The ionization energy is the energy necessary to
remove an electron completely from an atom.
X --> X+1 + e-
Group Trend - Electrons going into orbitals with
larger principal energy levels(further from the
nucleus)
Period Trend - electrons going into orbitals with
same principle energy level
What else is going on across a period?
Atomic Radius
Radii are estimated from actual
spacing in metals or molecules
Ionic Radius
• Cations
• larger or smaller than original atom?
• Group Trend?
• Period Trend?
• Anions
• larger or smaller than original atom?
• Group Trend?
• Period Trend?
• Mixed comparison
22
Ionic Radius Trends
Ionic radius of most
common ion reported in
picometers.
The size typically
decrease across the
period with a large jump
when going from anion to
cation.
Also, cations are smaller
than their atoms and
anions are larger than
their atoms.
Electronegativity Trends
Electronegativity is the ability of an atom to attract electrons to
itself in a chemical bond. It generally increases across a
period and decreases down a group.